Chemical Bonding / Unit 3

Introduction:

The world would be rather boring if it wasn't for chemical bonding. Without bonding, the world would lack all of its variety and complexity because we wouldn't have compounds. The world would be limited to the approximately 100 elementary substances we see in the periodic table. However, there are millions of different compounds and new ones are discovered on a regular basis. Obviously, atoms would rather be combined with one another than exist by themselves. Atoms find peace, love, and happiness (not to mention stability) when they combine by forming chemical bonds. In this unit, we will study how and why atoms go about combining and what they look like after they form compounds. Chemical bonds are forces of attraction that hold atoms together in groups called compounds. When atoms come in contact with each other, it will be the outermost electrons which interact. In the previous unit, we defined these as valence electrons. Valence electrons have more energy and are less stable than the inner electrons and will participate in chemical reactions. We will learn how valence electrons interact to form chemical bonds.

Two Main Types of Bonding:

I. Ionic Bonding - involves a transfer of electrons from one atom to another. Creates ions of opposite charge that experience an electrostatic attraction.

  • The forces of attraction between oppositely charged ions in the compound MN are called ionic bonds.
  • ion formation is most likely to occur when one of the atoms involved gives up electrons readily and the other atom has a strong attraction for electrons.
  • atoms will form a compound only when they will become more stable in doing so. What must happen in terms of energy for the atoms to become more stable?

Properties of Ionic Compounds:

1. Ionic compounds form crystals

  • the simplest ratio of the crystal is called a “formula unit”. If an aluminum chloride crystal contains 3 times as many chloride ions as aluminum ions than we use a formula unit of AlCl3.
  • there is really no such thing as an NaCl “molecule” because no pair of ions can be singled out as one unit. Cations are surrounded by anions and vice versa so that repulsions between like charges are minimized and attractions between opposite charges are maximized stable

2. Ionic compounds are solids at room temperature with high melting points

  • to melt an ionic solid, all bonds must be broken and this requires a high temp. because of the strong electrical forces that hold the crystal together.
  • only at high temps. will the oppositely charged ions have enough kinetic energy to break away from one another. MP of NaCl = 800 *C

3. Ionic Compounds are good conductors of electricity

  • must be melted (molten) or dissolved in solution (aq) so that the ions are free to move and are thus able to carry electrical charge.

Hydrogen gas (H2) is a diatomic molecule with a bond between two identical atoms. Which hydrogen atom will want to give up an electron and which one will want to receive an electron? Neither! They will “share”.

II. Covalent Bonding–when two atoms share a pair of pairs of valence electrons

  • Since the formation of a covalent bond involves sharing electrons, it occurs when the two atoms involved have similar attractions for electrons.
  • covalent bonds occur between nonmetals.
  • all “molecules” are held together by covalent bonds and understanding covalent bonding will enable us to study molecular structure.
  • again, if atoms combine, what must be true?

How are the atoms in a covalent bond becoming more stable?

  • We will consider a simple molecule: H2
  • when 2 H atoms approach each other closely, their 1s orbitals overlap. The two electrons are now attracted to their own nucleus and that of the other atom so they now spend their time between the atoms.
  • both atoms have both electrons so they have acquired a noble gas configuration by sharing electrons (1s2 He)

Properties of Molecular Substances:

  • molecular substances are gases, liquids, and low-melting solids at normal conditions.
  • molecules tend to be weakly attracted to other molecules of the same kind.
  • 2 major differences between ionic compounds and molecular substances:

1.Molecular substances have much lower melting and boiling points. To boil or melt a molecular substance, we do not have to break chemical bonds, just separate the molecules. No covalent bonds have to be broken. This requires little energy compared to breaking bonds between oppositely charged ions.

2.Molecular substances do not conduct electricity in the pure state because they consist of uncharged molecules.

Lewis Dot Symbols/Structures - American chemist Gilbert Lewis, 1916

  • Lewis Dot Symbols are simply a way of drawing an atom to show how many valence electrons it holds. Lewis Dot Structures are one way to predict the bonding that takes place in a compound.

Octet Rule - in forming covalent bonds, atoms tend to acquire noble gas configurations.

Writing Lewis structures/electron dot diagrams for molecules:

1.Using the formula of the molecule count the total number of valence electrons supplied by the atoms involved.

2.Draw the skeleton of the molecule

3.Distribute the electrons (dots) to give each atom an octet (except H) with all single bonds.

4.Count electrons used and adjust as needed.

*Some examples of exceptions to the octet rule: NO, BF3, PCl5, SF6, etc.

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Chemistry I Cary Academy W.G. Rushin